1. Chemistry Course CHE101: Chemistry Electrochemistry AP
Department of Chemistry
Lovely Professional University, Phagwara, Punjab, India.
Electrochemistry
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2. Electrochemistry
Chapter 19
Raymond Chang
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3. Oxidation half-reaction (lose e-)
Electrochemical processes are oxidation-reduction reactions in which:
the energy released by a spontaneous reaction is converted to electricity or
electrical energy is used to cause a nonspontaneous reaction to occur 2+ 2-
2Mg (s) + O2 (g) MgO (s)
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O2 + 4e O2-
Reduction half-reaction (gain e-)
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4. Terminology for Redox Reactions
OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen.
REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.
OXIDIZING AGENT—electron acceptor; species is reduced.
REDUCING AGENT—electron donor; species is oxidized.
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5. You can’t have one… without the other!
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons.
You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation
LEO the lion says GER!
GER!
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6. Another way to remember
OIL RIG
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7. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1.
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8. HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4
The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.
Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
HCO3-
Oxidation numbers of all the atoms in HCO3- ?
O = -2
H = +1
3x(-2) ? = -1
C = +4
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9. Balancing Redox Equations
The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution?
Write the unbalanced equation for the reaction ion ionic form.
Fe2+ + Cr2O Fe3+ + Cr3+
Separate the equation into two half-reactions.
Fe Fe3+ +2 +3
Oxidation:
Cr2O Cr3+ +6 +3
Reduction:
Balance the atoms other than O and H in each half-reaction.
Cr2O Cr3+
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10. Balancing Redox Equations
For reactions in acid, add H2O to balance O atoms and H+ to balance H atoms.
Cr2O Cr3+ + 7H2O
14H+ + Cr2O Cr3+ + 7H2O
Add electrons to one side of each half-reaction to balance the charges on the half-reaction.
Fe Fe3+ + 1e-
6e- + 14H+ + Cr2O Cr3+ + 7H2O
If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients.
6Fe Fe3+ + 6e-
6e- + 14H+ + Cr2O Cr3+ + 7H2O
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11. Balancing Redox Equations
Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel.
Oxidation:
6Fe Fe3+ + 6e-
Reduction:
6e- + 14H+ + Cr2O Cr3+ + 7H2O
14H+ + Cr2O Fe Fe3+ + 2Cr3+ + 7H2O
Verify that the number of atoms and the charges are balanced.
14x1 – 2 + 6x2 = 24 = 6x3 + 2x3
For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation.
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12. Galvanic Cells anode oxidation cathode reduction spontaneous
redox reaction
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13. Galvanic Cells
The difference in electrical potential between the anode and cathode is called:
cell voltage
electromotive force (emf)
cell potential
Cell Diagram
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode
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14. Standard Reduction Potentials
Standard reduction potential (E0) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M and all gases are at 1 atm.
Reduction Reaction
2e- + 2H+ (1 M) H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)
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15. Standard Reduction Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s) Zn2+ (1 M) + 2e-
Cathode (reduction):
2e- + 2H+ (1 M) H2 (1 atm)
Zn (s) + 2H+ (1 M) Zn2+ + H2 (1 atm)
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16. Standard Reduction Potentials
E0 = 0.76 V
cell
Standard emf (E0 )
cell
E0 = Ecathode - Eanode
cell
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
E0 = EH /H - EZn /Zn
cell + 2+ 2
0.76 V = 0 - EZn /Zn 2+
EZn /Zn = V 2+
Zn2+ (1 M) + 2e Zn E0 = V
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17. Standard Reduction Potentials
E0 = 0.34 V
cell
E0 = Ecathode - Eanode
cell
Ecell = ECu /Cu – EH /H 2+ + 2
0.34 = ECu /Cu - 0 2+
ECu /Cu = 0.34 V 2+
Pt (s) | H2 (1 atm) | H+ (1 M) || Cu2+ (1 M) | Cu (s)
Anode (oxidation):
H2 (1 atm) H+ (1 M) + 2e-
Cathode (reduction):
2e- + Cu2+ (1 M) Cu (s)
H2 (1 atm) + Cu2+ (1 M) Cu (s) + 2H+ (1 M)
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18. E0 is for the reaction as written
The more positive E0 the greater the tendency for the substance to be reduced
The half-cell reactions are reversible
The sign of E0 changes when the reaction is reversed
Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E0
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19. Electrochemical series
Arrangement of elements in the order of increasing reduction potential
Applications:
1. Relative strengths of oxidising and reducing agents: down-decreasing tendency to get reduced/low reduction potential/oxidising agent-weak
Subs with low reduction potential are stronger reducing agents and vice versa

20. F2 (weakest reducing agent)strongest oxidizing agent because it has the greatest tendency to be reduced.
Li (strongest reducing agent)is the weakest oxidizing agent because it is the most difficult species to reduce
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21. Cd is the stronger oxidizer
Applications contd..
2. Calculation of EMF of the cell
What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e Cd (s) E0 = V
Cd is the stronger oxidizer
Cd will oxidize Cr
Cr3+ (aq) + 3e Cr (s) E0 = V
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22. Cd is the stronger oxidizer
What is the standard emf of an electrochemical cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a Cr electrode in a 1.0 M Cr(NO3)3 solution?
Cd2+ (aq) + 2e Cd (s) E0 = V
Cd is the stronger oxidizer
Cd will oxidize Cr
Cr3+ (aq) + 3e Cr (s) E0 = V
Anode (oxidation):
Cr (s) Cr3+ (1 M) + 3e-
x 2
Cathode (reduction):
2e- + Cd2+ (1 M) Cd (s)
x 3
2Cr (s) + 3Cd2+ (1 M) Cd (s) + 2Cr3+ (1 M)
E0 = Ecathode - Eanode
cell
E0 = – (-0.74)
cell
E0 = 0.34 V
cell
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23. Applications contd…
3. Predicting the feasibility of the reaction: species to lose electron must have lower reduction potential
E˚: Cu=+0.34V and Ag=+0.80V Ag has greater tendency to get reduced.
Cu(s) + 2Ag+(aq)Cu2+(aq)+2Ag(s)
Is feasible while reverse is not
NOTE: E˚ sign changes but its magnitude remains the same when we reverse a reaction
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24. The half-cell reactions are reversible
Depending on the conditions, any electrode can act either as an anode or as a cathode.(SHE is the cathode (H is reduced to H2) when coupled with zinc in a cell and that it becomes the anode (H2 is oxidized to H) when used in a cell with copper.)
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25. Zn spontaneously reduces Cu2+ to form Zn2+ and Cu.
Under standard-state conditions, any species on the left of a given half-cell reaction will react spontaneously with a species that appears on the right of any half-cell reaction located above it in Table. This principle is sometimes called the diagonal rule. In the case of the Cu/Zn electrochemical cell
Zn spontaneously reduces Cu2+ to form Zn2+ and Cu.
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26. electrode potentials are intensive properties
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27. Applications contd…
4.To predict whether a metal can liberate hydrogen from acid or not: M(E˚=-ive)
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30. Spontaneity of Redox Reactions
DG = -nFEcell
n = number of moles of electrons in reaction
F = 96,500 J
V • mol
DG0 = -nFEcell
= 96,500 C/mol
DG0 = -RT ln K
= -nFEcell
Ecell = RT nF
ln K
(8.314 J/K•mol)(298 K)
n (96,500 J/V•mol)
ln K = = V n
ln K
Ecell = V n
log K
Ecell
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31. Spontaneity of Redox Reactions
DG0 = -RT ln K
= -nFEcell
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32. What is the equilibrium constant for the following reaction at 250C
What is the equilibrium constant for the following reaction at 250C? Fe2+ (aq) + 2Ag (s) Fe (s) + 2Ag+ (aq) = V n
ln K
Ecell
Oxidation:
2Ag Ag+ + 2e-
n = 2
Reduction:
2e- + Fe Fe
E0 = EFe /Fe – EAg /Ag 2+ +
E0 = – (0.80)
E0 = V V
x n E0
cell
exp
K = V
x 2
-1.24 V
= exp
K = 1.23 x 10-42
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33. The Effect of Concentration on Cell Emf
DG = DG0 + RT ln Q
DG = -nFE
DG0 = -nFE
-nFE = -nFE0 + RT ln Q
Nernst equation
E = E0 -
ln Q RT nF
At 298 - V n
ln Q E
E = - V n
log Q E
E =
At equilibrium, there is no net transfer of electrons, so E =0 and Q= K, where K is the equilibrium constant.(Q= reaction quotient)
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34. Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
Will the following reaction occur spontaneously at 250C if [Fe2+] = 0.60 M and [Cd2+] = M?
Fe2+ (aq) + Cd (s) Fe (s) + Cd2+ (aq)
Oxidation:
Cd Cd2+ + 2e-
n = 2
Reduction:
2e- + Fe Fe
E0 = EFe /Fe – ECd /Cd 2+
E0 = – (-0.40) - V n
ln Q E
E =
E0 = V - V 2 ln
-0.04 V
E =
0.010
0.60
E = 0.013
E > 0
Spontaneous
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35. Nickel cadmium storage cell
Batteries: A battery is an electrochemical cell, or a series of combined electrochemical cells, that can be used as a source of direct electric current at a constant voltage.
Primary cells
Rxns cant be reversed.
Not rechargable
Dry cell
Mercury cell
Secondary Cells
Rxns can be reversed.
Rechargable
Lead storage battery
Nickel cadmium storage cell
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36. Dry cell Leclanché cell Anode: Zn (s) Zn2+ (aq) + 2e- Cathode:
a cell without a fluid component
Anode:
Zn (s)  Zn2+ (aq)
Zn (s) Zn2+ (aq) + 2e-
NH4+ (aq) + MnO2 (s) Mn2O3 (s) + NH3 (aq)
Cathode:
2NH4 (aq) + 2MnO2 (s) + 2e Mn2O3 (s) + 2NH3 (aq) + H2O (l) +
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
NH3 is not liberated as gas but combines with Zn as [Zn(NH3)2]2+
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38. Dry Cell Disadvantage Uses
During use Zn gets consumed. In the end holes get developed which are responsible for leakage.
Overcome: iron/steel sheet covering the Zn
NH4Cl being acidic corrodes Zn container even when not in use
Torches, toys, flash lights, calculators, tape recorders etc.
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39. Batteries Mercury Battery Anode:
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-
Cathode:
HgO (s) + H2O (l) + 2e Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
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41. Mercury Cell Advantages Uses
Cell overall rxn doesn’t involve any ion whose conc. can be changed(no change in electrolyte composition during operation)
Potential remains constant (1.35 V) contrary to dry cell whose potential decreases with time
It also has a considerably higher capacity and longer life.
Cameras, watches, pacemakers, hearing aids, electric watches, and light meters.
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42. Batteries Lead storage battery Anode:
Pb (s) + SO2- (aq) PbSO4 (s) + 2e- 4
Cathode:
PbO2 (s) + 4H+ (aq) + SO2- (aq) + 2e PbSO4 (s) + 2H2O (l) 4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO2- (aq) PbSO4 (s) + 2H2O (l) 4
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44. Solid State Lithium Battery
Anode- lithium in graphite
Graphite will be having tiny pores which can hold lithium and lithium ions
Cathode--- transition metal oxide such as co02 which can also hold lithium ions
Electrolyte-non aq i.e organic solvent and dissolvd salt
Solid State Lithium Battery

45. The advantage in choosing lithium as the anode is that it has the most negative E° value (-3.04V).
Furthermore, lithium is a lightweight metal, so that only g of Li (its molar mass) are needed to produce 1 mole of electrons.
The electrolyte is a polymer (poly acrylonitrile or polyethyllene oxide) material that permits the passage of ions but not of electrons.
The cathode is made of either TiS2 or V6O13.
The cell voltage of a solid-state lithium battery can be as high as 3 V, and it can be recharged in the same way as a lead storage battery.

46. CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
Fuel cells
POWER PLANTS
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
Only 40% efficiency
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47. Batteries
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning
Anode:
2H2 (g) + 4OH- (aq) H2O (l) + 4e-
Cathode:
O2 (g) + 2H2O (l) + 4e OH- (aq)
2H2 (g) + O2 (g) H2O (l)
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48. The standard emf of the cell can be calculated as follows
E° = E°ox - E°red
= 0.83 V – (-0.40 V)
= 1.23 V
The reaction is the same as the hydrogen combustion reaction, but the oxidation and reduction are carried out separately at the anode and the cathode.
Like platinum in the standard hydrogen electrode, the electrodes have a twofold function. They serve as electrical conductors, and they provide the necessary surfaces for the initial decomposition of the molecules into atomic species, prior to electron transfer. They are electrocatalysts.
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49. Unlike batteries, fuel cells do not store chemical energy.
Disadvantages:
Unlike batteries, fuel cells do not store chemical energy.
Reactants must be constantly resupplied, and products must be constantly removed from a fuel cell.
In this respect, a fuel cell resembles an engine more than it does a battery.
Advantages
fuel cells may be as much as 70 percent efficient, about twice as efficient as an internal combustion engine.
In addition, fuel-cell generators are free of the noise, vibration, heat transfer, thermal pollution, and other problems normally associated with conventional power plants.
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50. Corrosion
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51. Methods of prevention:
1. Painting:
If the paint is scratched, pitted, or dented to expose even the smallest area of bare metal, rust will form under the paint layer.
2. Passivation:
The surface of iron metal can be made inactive by a process called passivation.
A thin oxide layer is formed when the metal is treated with a strong oxidizing agent such as concentrated nitric acid.
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52. 3. Coating with other metals:
An iron container can be covered with a layer of another metal such as tin or zinc.
A “tin” can is made by applying a thin layer of tin over iron.
Rust formation is prevented as long as the tin layer remains intact.
However, once the surface has been scratched, rusting occurs rapidly.
Fe(s)  Fe2+ + 2e- E° = V
Sn2+(aq) + 2e-  Sn(s) E° = 0.14 V
2. The protective process is different for zinc-plated, or galvanized, iron.
Zinc is more easily oxidized than iron
Zn(s)  Zn2+(aq) + 2e- E° = 0.76 V
So even if a scratch exposes the iron, the zinc is still attacked.
In this case, the zinc metal serves as the anode and the iron is the cathode.

53. Cathodic Protection of an Iron Storage Tank
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